I'm watching some (really awesome) chemistry courses, and I think I have a fairly decent handle on acids and bases (or maybe not, and I'm fooling myself). But the teacher just showed a table of weak/strong acids. $\ce{HSO4-}$ was one of them.

I'm wondering, why hydrogen sulfate ion instead of just hydrogen sulfate? Why is $\ce{HSO4-}$ an acid and $\ce{HSO4}$ apparently isn't? This was the only ion on the table. Or does $\ce{HSO4}$ simply not exist?


3 Answers 3


The bisulfate molecule doesn't exist in solution. I can only find references to the hydrogen sulfate molecule as part of salts, which implies that hydrogen sulfate is usually found as an ion.

The bisulfate (hydrogen sulfate) ion is indeed an acid. It is a relatively strong weak acid too, with a $\ce{K_a}$ value of $1.2 * 10^{-2}$.

Part of the reason for its acidity has to do with its electronegative oxygens isolationg electron density away from the hydrogen atom bonded to the oxygen. This makes the hydrogen more partially positive and thus more reactive (more likely to be taken by a base).

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  • $\begingroup$ I have to add, that while $\ce{HSO4}$ does not exist, the doubled molecule $\ce{H2S2O8}$ (peroxodisulfuric acid) does exist and is apparently a quite strong acid (its double ammonium salt is stable solid) $\endgroup$
    – permeakra
    Commented Jul 18, 2014 at 16:54
  • 1
    $\begingroup$ I have to add, that $\ce{HSO4}$ might well exist, just not in aqueous solution - in the gas phase, why not ? $\endgroup$ Commented Jul 18, 2014 at 18:32
  • $\begingroup$ I agree Martin! $\endgroup$
    – Dissenter
    Commented Jul 19, 2014 at 6:57
  • $\begingroup$ @Martin - it would have an odd number of valence electrons, and so it probably wouldn't exist for long! $\endgroup$
    – thomij
    Commented Jul 25, 2014 at 19:00
  • $\begingroup$ "Part of the reason for its acidity has to do with its electronegative oxygens isolationg electron density away from the hydrogen atom bonded to the oxygen. " is that actually true? How would that work? Isn't it just acidic because the negative charge can be fairly good delocalized into the $\pi$-bonds? If we'd were to replace one of the two S-O double bond with a SOH group, I think it would be that acidic anymore... $\endgroup$
    – Jori
    Commented Jul 25, 2014 at 20:58

I notice that there is no mention of polyprotic acids in the other answers!

$\ce{H2SO4}$ (sulfuric acid) is diprotic - meaning it has two "detachable" protons which can come off in aqueous solution. The first one is strongly acidic - the $Ka = 2.4 × 10^6$ - this means the odds of finding a non-dissociated $\ce{H2SO4}$ molecule in solution are something like 1 out of 2.4 million.

On the other hand, the second proton is weakly acidic - the $Ka = 1.0 × 10^{-2}$. This means that there will mostly be $\ce{HSO4-}$ instead of $\ce{SO4^2-}$ in solution (roughly 100 times as much).

The other answers are correct in saying that $\ce{HSO4}$ does not exist in solution - but it is important to realize that we list the bisulfate ion ($\ce{HSO4-}$) as a weak acid because it is the conjugate base of sulfuric acid, and is in turn the conjugate acid of the sulfate ion. This becomes very important when you start to study acid/base equilibria.


$\ce{HSO4}$ does not exist. There is $\ce{H2SO4}$ and $\ce{HSO4-}$ Try drawing the Lewis structure for sulfate ion. You'll see that there are two oxygens with formal charges. They can be left as is or they can have $\ce{H}$ bonded to them.

  • $\begingroup$ I think you mean H- bonded to them (H forms single bonds). $\endgroup$
    – Dissenter
    Commented Jul 18, 2014 at 16:32
  • $\begingroup$ The adequate Lewis structure for the sulfate ion has one negative charge at each oxygen and a double positive charge at sulfur. Everything else is just too far from the truth. $\endgroup$ Commented Jul 18, 2014 at 18:34

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